Chapter Fifteen 1 Acids Bases and AcidBase Equilibria

2 The Brønsted–Lowry Theory • Arrhenius theory: an acid forms H+ in water; and a base forms OH– in water. • But not all acid–base reactions involve water, and many bases (NH 3, carbonates) do not contain any OH–. • Brønsted–Lowry theory defines acids and bases in terms of proton (H+) transfer. • A Brønsted–Lowry acid is a proton donor. • A Brønsted–Lowry base is a proton acceptor. • The conjugate base of an acid is the acid minus the proton it has donated. • The conjugate acid of a base is the base plus the accepted proton. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

3 Ionization of HCl acts as an acid by donating H+ to H 2 O is a base in this reaction because it accepts the H+ Conjugate acid of H 2 O Conjugate base of HCl Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

4 Ionization of Ammonia Prentice Hall © 2005 General Chemistry 4 th edition, Hill,

5 Water Is Amphiprotic H 2 O acts as an acid when it donates H+, forming the conjugate base ___ H 2 O acts as a base when it accepts H+, forming the conjugate acid ___ Amphiprotic: Can act as either an acid or as a base Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

6 Example 15. 1 Identify the Brønsted–Lowry acids and bases and their conjugates in: (a) H 2 S + NH 3 NH 4+ + HS– (b) OH– + H 2 PO 4– H 2 O + HPO 42– Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

7 Ka and Kb The equilibrium constant for a Brønsted acid is represented by Ka, and that for a base is represented by Kb. CH 3 COOH(aq) + H 2 O(l) H 3 O+(aq) + CH 3 COO–(aq) [H 3 O+][CH 3 COO–] Ka = ––––––––– [CH 3 COOH] Notice that H 2 O is not included in either equilibrium expression. NH 3(aq) + H 2 O(l) NH 4+(aq) + OH–(aq) [NH 4+][OH–] Kb = ––––––– [NH 3] Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

8 Strength of Conjugate Acid–Base Pairs • A stronger acid can donate H+ more readily than a weaker acid. • The stronger an acid, the weaker is its conjugate base. • The stronger a base, the weaker is its conjugate acid. • An acid–base reaction is favored in the direction from the stronger member to the weaker member of each conjugate acid–base pair. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

9 … the weaker the conjugate base. The stronger the acid … And the stronger the base … … the weaker the conjugate acid. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

10 Acid/Base Strength and Direction of Equilibrium • In Table 15. 1, HBr lies above CH 3 COOH in the acid column. • Since HBr is a stronger acid than CH 3 COOH, the equilibrium for the reaction: Weaker base Weaker acid Stronger base Stronger acid lies to the left. • We reach the same conclusion by comparing the strengths of the bases (right column of Table 15. 1). • CH 3 COO– lies below Br– ; CH 3 COO– is the stronger base: Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

11 Strong Acids • The “strong” acids—HCl, HBr, HI, HNO 3, H 2 SO 4, HCl. O 4—are considered “strong” because they ionize completely in water. • The “strong” acids all appear above H 3 O+ in Table 15. 1. • The strong acids are leveled to the same strength—to that of H 3 O+—when they are placed in water. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

12 Periodic Trends in Acid Strength • The greater the tendency for HX (general acid) to transfer a proton to H 2 O, the more the forward reaction is favored and the stronger the acid. • A factor that makes it easier for the H+ to leave will increase the strength of the acid. • Acid strength is inversely proportional to H—X bond-dissociation energy. Weaker H—X bond => stronger acid. • Acid strength is directly proportional to anion radius. Larger X radius => stronger acid. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

13 Periodic Trends in Acid Strength Prentice Hall © 2005 General Chemistry 4 th

14 Strength of Oxoacids Acid strength increases with the electronegativity of the central atom, and with the number of terminal oxygen atoms. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

15 Strength of Carboxylic Acids • Carboxylic acids all have the –COOH group in common. • Differences in acid strength come from differences in the R group attached to the carboxyl group. • In general, the more that electronegative atoms appear in the R group, the stronger is the acid. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

16 Example 15. 2 Select the stronger acid in each pair: (a) nitrous acid, HNO 2, and nitric acid, HNO 3 (b) Cl 3 CCOOH and Br. CH 2 COOH Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

17 Strengths of Amines as Bases • Aromatic amines are much weaker bases than aliphatic amines. • This is due in part to the fact that the p electrons in the benzene ring of an aromatic molecule are delocalized and can involve the nitrogen atom’s lone-pair electrons in the resonance hybrid. • As a result, the lone-pair electrons are much less likely to accept a proton. • Electron-withdrawing groups on the ring further diminish the basicity of aromatic amines relative to aniline. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

18 Example 15. 3 Select the weaker base in each pair: Prentice Hall ©

19 Self-Ionization of Water • Even pure water conducts some electricity. This is due to the fact that water self-ionizes: • The equilibrium constant for this process is called the ion product of water (Kw). • At 25 °C, Kw = 1. 0 x 10– 14 = [H 3 O+][OH–] • This equilibrium constant is very important because it applies to all aqueous solutions—acids, bases, salts, and nonelectrolytes—not just to pure water. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

20 The p. H Scale • Concentration of H 3 O+ can vary over a wide range in aqueous solution, from about 10 M to about 10– 14 M. • A more convenient expression for H 3 O+ is p. H = –log [H 3 O+] and so [H 3 O+] = 10–p. H • The “negative logarithm” function of p. H is so useful that it has been applied to other species and constants. p. OH = –log [OH–] and so [OH–] = 10–p. OH p. Kw = –log Kw • At 25 °C, p. Kw = 14. 00 • p. Kw = p. H + p. OH = 14. 00 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

21 The p. H Scale Since p. H is a logarithmic scale, cola drinks (p. H about 2. 5) are about ____ times as acidic as tomatoes (p. H about 4. 5) Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

22 Example 15. 4 By the method suggested in Figure 15. 5, a student determines the p. H of milk of magnesia, a suspension of solid magnesium hydroxide in its saturated aqueous solution, and obtains a value of 10. 52. What is the molarity of Mg(OH)2 in its saturated aqueous solution? The suspended, undissolved Mg(OH)2(s) does not affect the measurement. Example 15. 5 A Conceptual Example Is the solution 1. 0 x 10– 8 M HCl acidic, basic, or neutral? Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

23 Equilibrium in Solutions of Weak Acids and Weak Bases 1. 2. 3. 4. 5. These calculations are similar to the equilibrium calculations performed in Chapter 14. An equation is written for the reversible reaction. Data are organized, often in an ICE format. Changes that occur in establishing equilibrium are assessed. Simplifying assumptions are examined (the “ 5% rule”). Equilibrium concentrations, equilibrium constant, etc. are calculated. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

24 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary,

25 Example 15. 6 Ordinary vinegar is approximately 1 M CH 3 COOH and as shown in Figure 15. 6, it has a p. H of about 2. 4. Calculate the expected p. H of 1. 00 M CH 3 COOH(aq), and show that the calculated and measured p. H values are in good agreement. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

26 Example 15. 7 What is the p. H of 0. 00200 M Cl.

27 Example 15. 8 What is the p. H of 0. 500 M NH 3(aq)? Example 15. 9 The p. H of a 0. 164 M aqueous solution of dimethylamine is 11. 98. What are the values of Kb and p. Kb? The ionization equation is (CH 3)2 NH + H 2 O (CH 3)2 NH 2+ + OH– Kb = ? Dimethylamine Dimethylammonium ion Example 15. 10 A Conceptual Example Without doing detailed calculations, indicate which solution has the greater [H 3 O+], 0. 030 M HCl or 0. 050 M CH 3 COOH. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

28 Polyprotic Acids • A monoprotic acid has one ionizable H atom per molecule. • A polyprotic acid has more than one ionizable H atom per molecule. – Sulfuric acid, H 2 SO 4 – Carbonic acid, H 2 CO 3 – Phosphoric acid, H 3 PO 4 Diprotic Triprotic • The protons of a polyprotic acid dissociate in steps, each step having a value of Ka. • Values of Ka decrease successively for a given polyprotic acid. Ka 1 > Ka 2 > Ka 3 , etc. • Simplifying assumptions may be made in determining the concentration of various species from polyprotic acids. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

29 Example 15. 11 Calculate the following concentrations in an aqueous solution that is 5. 0 M H 3 PO 4: (a) [H 3 O+] (b) [H 2 PO 4–] (c) [HPO 42–] (d) [PO 43–] Example 15. 12 What is the approximate p. H of 0. 71 M H 2 SO 4? Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

30 Ions as Acids and Bases • HCl is a strong acid, therefore Cl– is so weakly basic in water that a solution of chloride ions (such as Na. Cl) is virtually neutral. • Acetic acid, CH 3 COOH, is a weak acid, so acetate ion, CH 3 COO–, is significantly basic in water. • A solution of sodium acetate (which dissociates completely into sodium and acetate ions in water) is therefore slightly basic: CH 3 COO– + H 2 O CH 3 COOH + OH– Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

31 Carbonate Ion as a Base A carbonate ion accepts a proton from water, leaving behind an OH– and making the solution basic. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

32 Ions as Acids and Bases (cont’d) • Salts of strong acids and strong bases form neutral solutions: Na. Cl, KNO 3 • Salts of weak acids and strong bases form basic solutions: KNO 2, Na. Cl. O • Salts of strong acids and weak bases form acidic solutions: NH 4 NO 3 • Salts of weak acids and weak bases form solutions that may be acidic, neutral, or basic; it depends on the relative strengths of the cations and the anions: NH 4 NO 2, CH 3 COONH 4. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

33 Example 15. 13 A Conceptual Example (a) Is NH 4 I(aq) acidic, basic, or neutral? (b) What conclusion can you draw from Figure 15. 8 d about the equilibrium constants for the hydrolysis reactions in CH 3 COONH 4(aq)? Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

34 Ions as Acids and Bases (cont’d) • In order to make quantitative predictions of p. H of a salt solution, we need an equilibrium constant for the hydrolysis reaction. • The relationship between Ka and Kb of a conjugate acid–base pair is: Ka x Kb = Kw • If instead we have values of p. Ka or p. Kb: p. Ka + p. Kb = p. Kw = 14. 00 (at 25 °C) Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

35 Example 15. 14 Calculate the p. H of a solution that is 0. 25 M CH 3 COONa(aq). Example 15. 15 What is the molarity of an NH 4 NO 3(aq) solution that has a p. H = 4. 80? Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

36 The Common Ion Effect • Consider a solution of acetic acid. • If we add acetate ion as a second solute (i. e. , sodium acetate), the p. H of the solution increases: Le. Châtelier’s principle: What happens to [H 3 O+] when the equilibrium shifts to the left? Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

37 The Common Ion Effect (cont’d) The common ion effect is the suppression of the ionization of a weak acid or a weak base by the presence of a common ion from a strong electrolyte. Acetic acid solution at equilibrium: a few H 3 O+ ions and a few CH 3 COO– ions Prentice Hall © 2005 When acetate ion is added, and equilibrium reestablished: more acetate ions, fewer H 3 O+ ions General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

38 Example 15. 16 Calculate the p. H of an aqueous solution that is both 1. 00 M CH 3 COOH and 1. 00 M CH 3 COONa. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

39 Buffer Solutions • Many industrial and biochemical reactions—especially enzyme-catalyzed reactions—are sensitive to p. H. • To work with such reactions we often need a solution that maintains a nearly-constant p. H. • A buffer solution is a solution that changes p. H only slightly when small amounts of a strong acid or a strong base are added. • A buffer contains significant concentrations of both – a weak acid and its conjugate base, or – a weak base and its conjugate acid. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

40 Buffer Solutions (cont’d) • The acid component of the buffer neutralizes small added amounts of OH–, forming the weaker conjugate base which does not affect p. H much: HA + OH– H 2 O + A– • The base component neutralizes small added amounts of H 3 O+, forming the weaker conjugate acid which does not affect p. H much. A– + H 3 O+ H 2 O + HA • Pure water does not buffer at all … Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

Pure water increases in p. H by about 5 p. H units when the OH– is added, and decreases by about 5 p. H units when the H 3 O+ is added. 41 In contrast, the same amounts of OH– and H 3 O+ added to a buffer solution barely change the p. H. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

42 An Equation for Buffer Solutions • In certain applications, there is a need to repeat the calculations of the p. H of buffer solutions many times. This can be done with a single, simple equation, but there are some limitations. • The Henderson–Hasselbalch equation: [conjugate base] p. H = p. Ka + log ––––––– [weak acid] • To use this equation, the ratio [conjugate base]/[weak acid] must have a value between 0. 10– 10 and both concentrations must exceed Ka by a factor of 100 or more. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

43 Buffer Capacity and Buffer Range • There is a limit to the ability of a buffer solution to neutralize added acid or base. • This buffer capacity is reached before either buffer component has been consumed. • In general, the more concentrated the buffer components in a solution, the more added acid or base the solution can neutralize. • As a rule, a buffer is most effective if the concentrations of the buffer acid and its conjugate base are equal or nearly so. • Therefore, a buffer is most effective when the desired p. H of the buffer is very near p. Ka of the weak acid of the buffer. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

44 Example 15. 17 A buffer solution is 0. 24 M NH 3 and 0. 20 M NH 4 Cl. (a) What is the p. H of this buffer? (b) If 0. 0050 mol Na. OH is added to 0. 500 L of this solution, what will be the p. H? Example 15. 18 What concentration of acetate ion in 0. 500 M CH 3 COOH(aq) produces a buffer solution with p. H = 5. 00? Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

45 Acid–Base Indicators • An acid–base indicator is a weak acid or base. • The acid form (HA) of the indicator has one color, the conjugate base (A–) has a different color. One of the “colors” may be colorless. • In an acidic solution, [H 3 O+] is high. Because H 3 O+ is a common ion, it suppresses the ionization of the indicator acid, and we see the color of HA. • In a basic solution, [OH–] is high, and it reacts with HA, forming the color of A–. • Acid–base indicators are often used for applications in which a precise p. H reading isn’t necessary. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

Different indicators have different values of Ka, so they exhibit color changes at different values of p. H … Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry 46 Chapter Fifteen

47 Example 15. 19 A Conceptual Example Explain the series of color changes of thymol blue indicator produced by the actions pictured in Figure 15. 14: (a) A few drops of thymol blue are added to HCl(aq). (b) Some sodium acetate is added to solution (a). (c) A small quantity of sodium hydroxide is added to solution (b). (d) An additional, larger quantity of sodium hydroxide is added to solution (c). Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

48 Neutralization Reactions • At the equivalence point in an acid–base titration, the acid and base have been brought together in precise stoichiometric proportions. • The endpoint is the point in the titration at which the indicator changes color. • Ideally, the indicator is selected so that the endpoint and the equivalence point are very close together. • The endpoint and the equivalence point for a neutralization titration can be best matched by plotting a titration curve, a graph of p. H versus volume of titrant. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

Titration Curve, Strong Acid with Strong Base 49 Bromphenol blue, bromthymol blue, and phenolphthalein all change color at very nearly 20. 0 m. L At about what volume would we see a color change if we used methyl violet as the indicator? Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

50 Example 15. 20 Calculate the p. H at the following points in the titration of 20. 00 m. L of 0. 500 M HCl with 0. 500 M Na. OH: H 3 O+ + Cl– + Na+ + OH– Na+ + Cl– + 2 H 2 O (a) before the addition of any Na. OH (b) after the addition of 10. 00 m. L of 0. 500 M Na. OH (c) after the addition of 20. 00 m. L of 0. 500 M Na. OH (d) after the addition of 20. 20 m. L of 0. 500 M Na. OH Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

Titration Curve, Weak Acid with Strong Base 51 The equivalence-point p. H is NOT 7. 00 here. Why not? ? Bromphenol blue was ok for the strong acid/strong base titration, but it changes color far too early to be useful here. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

52 Example 15. 21 Calculate the p. H at the following points in the titration of 20. 00 m. L of 0. 500 M CH 3 COOH with 0. 500 M Na. OH: CH 3 COOH + Na+ + OH– Na+ + CH 3 COO– + H 2 O (a) before the addition of any Na. OH (b) after the addition of 8. 00 m. L of 0. 500 M Na. OH (c) after the addition of 10. 00 m. L of 0. 500 M Na. OH (d) after the addition of 20. 00 m. L of 0. 100 M Na. OH (e) after the addition of 21. 00 m. L of 0. 100 M Na. OH Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

53 Example 15. 22 A Conceptual Example This titration curve shown in Figure 15. 18 involves 1. 0 M solutions of an acid and a base. Identify the type of titration it represents. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

54 Lewis Acids and Bases • There are reactions in nonaqueous solvents, in the gaseous state, and even in the solid state that can be considered acid –base reactions which Brønsted–Lowry theory is not adequate to explain. • A Lewis acid is a species that is an electron-pair acceptor and a Lewis base is a species that is an electron-pair donor. Ca. O(s) + SO 2(g) Ca. SO 3(s) Sulfur accepts an electron pair from the oxygen of Ca. O • In organic chemistry, Lewis acids are often called electrophiles (“electron-loving”) and Lewis bases are often called nucleophiles (“nucleus-loving”). Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen

55 Cumulative Example A 0. 0500 M aqueous solution of cyanoacetic acid, CNCH 2 COOH, has a freezing point of – 0. 11 °C. Calculate the freezing point of a 0. 250 M aqueous solution of cyanoacetic acid. Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, Mc. Creary, Perry Chapter Fifteen